Pathogenic microbes, molds, mildew, spores, and organic and inorganic pollutants are commonly found in the environment. Microbial control and disinfection in environmental spaces is desirable to improve health. Numerous ways have been used in the past in an attempt to purify air and disinfect surfaces. For example, it is already known that Reactive Oxygen Species (ROS) produced by, e.g., photocatalytic oxidation process can oxidize organic pollutants and kill microorganisms. More particularly, hydroxyl radical, hydroperoxyl radicals, chlorine and ozone, end products of the photocatalytic reaction, have been known to be capable of oxidizing organic compounds and killing microorganisms. However, there are limitations to the known methods and devices, not only due to efficacy limitation but also due to safety issues.
ROS is the term used to describe the highly activated air that results from exposure of ambient humid air to ultraviolet light. Light in the ultraviolet range emits photons at a frequency that when absorbed has sufficient energy to break chemical bonds. UV light at wavelengths of 250-255 nm is routinely used as a biocide. Light below about 181 nm, up to 182-187 nm is competitive with corona discharge in its ability to produce ozone. Ozonation and UV radiation are both being used for disinfection in community water systems. Ozone is currently being used to treat industrial wastewater and cooling towers.
Hydrogen peroxide is generally known to have antimicrobial properties and has been used in aqueous solution for disinfection and microbial control. Attempts to use hydrogen peroxide in the gas phase however, have previously been hampered by technical hurdles to the production of Purified Hydrogen Peroxide Gas (PHPG). Vaporized aqueous solutions of hydrogen peroxide produce an aerosol of microdroplets composed of aqueous hydrogen peroxide solution. Various processes for “drying” vaporized hydrogen peroxide (VHP) solutions produce, at best, a hydrated form of hydrogen peroxide. These hydrated hydrogen peroxide molecules are surrounded by water molecules bonded by electrostatic attraction and London Forces. Thus, the ability of the hydrogen peroxide molecules to directly interact with the environment by electrostatic means is greatly attenuated by the bonded molecular water, which effectively alters the fundamental electrostatic configuration of the encapsulated hydrogen peroxide molecule. Further, the lowest concentration of vaporized hydrogen peroxide that can be achieved is generally well above the 1.0 ppm Occupational Safety and Health Administration (OSHA) workplace safety limit, making these processes unsuitable for use in occupied areas.
Photocatalysts that have been demonstrated for the destruction of organic pollutants in fluid include but are not limited to TiO2, ZnO, SnO2, WO3, CdS, ZrO2, SB2O4, and Fe2O3. Titanium dioxide is chemically stable, has a suitable bandgap for UV/Visible photoactivation, and is relatively inexpensive. Therefore, photocatalytic chemistry of titanium dioxide has been extensively studied over the last thirty years for removal of organic and inorganic compounds from contaminated air and water.
Because photocatalysts can generate hydroxyl radicals from adsorbed water when activated by ultraviolet light of sufficient energy, they show promise for use in the production of PHPG for release into the environment when applied in the gas phase. Existing applications of photocatalysis, however, have focused on the generation of a plasma containing many different reactive chemical species. Further, the majority of the chemical species in the photocatalytic plasma are reactive with hydrogen peroxide, and inhibit the production of hydrogen peroxide gas by means of reactions that destroy hydrogen peroxide. Also, any organic gases that are introduced into the plasma inhibit hydrogen peroxide production both by direct reaction with hydrogen peroxide and by the reaction of their oxidized products with hydrogen peroxide.
The photocatalytic reactor itself also limits the production of PHPG for release into the environment. Because hydrogen peroxide has greater chemical potential than oxygen to be reduced as a sacrificial oxidant, it is preferentially reduced as it moves downstream in photocatalytic reactors as rapidly as it is produced by the oxidation of water.
TABLE 1Oxidation/Reduction Half ReactionsStandard ReductionPhoto-Activation of CatalystPotential (eV)hv ⇄ h+ + e− (on TiO2 catalyst)≤−3.2hv ⇄ h+ + e− (on TiO2 catalyst with co-catalyst)≤−2.85Loss of Free Electrons Due to Electron-HoleRecombinationh+ + e− ⇄ heat (on TiO2 catalyst)≥3.2h+ + e− ⇄ heat (on TiO2 catalyst with co-catalyst)≥2.85Formation of Hydroxyl Radicals (only if water isadsorbed on active sites on catalyst, preventingElectron-Hole Recombination)h+ + H2O ⇄ OH* + H+2.85Thermodynamically Favored loss of HydroxylRadicals by Free Electron Reduction in aConcentrated Plasma Reactor, but Avoided in aPHPG ReactorOH* + e− + H+⇄ H2O2.02Combination of Hydroxyl Radicals to FormHydrogen Peroxide is not ThermodynamicallyFavored Compared to Free Electron Reductionin a Plasma Reactor, but is Promoted by a PHPGReactor by Creating a Dilute Hydroxyl RadicalField Separated from Free Electrons2OH* ⇄ H2O21.77Spontaneous Reactions That would Destroy anyHydrogen Peroxide in Concentrated PlasmaReactors, but which are Avoided by a PHPGReactor by Creating a Dilute Hydroxyl RadicalField Separated from Free Electrons and Light2OH* + H2O2 ⇄ 2H2O + O22.805H2O2 + 2H+ + 2e− ⇄ 2H2O1.78H2O2 + hv ⇄ 2OH* (by Photolysis)1.77e− + H2O2 ⇄ OH* + OH−0.71Reactions that Create Hydrogen Peroxidethrough the Forced Reduction of Dioxygen in aPHPG Reactor, but not in a Concentrated PlasmaReactore− + O2 ⇄ O2− (First Step is non-Spontaneous)−0.132H+ + 2e− + O2 ⇄ H2O2 (Overall Reaction)0.70Other Reactions Common in a ConcentratedPlasma Reactor, but which do not take place in aPHPG Reactor, which does not use Ozone-Producing Wavelengths of LightO2 + hv ⇄ 2O* (by Photolysis)≤−5.132O* + 2O2 ⇄ 2O32.99O3 + 2H+ + 2e− ⇄ O2(g) + H2O2.075O3 + H2O + 2e− ⇄ O2(g) + 2OH−1.24Ozone Destruction of Hydrogen PeroxideO3 + H2O2 ⇄ H2O + 2O21.381
Additionally, several side reactions generate a variety of species that become part of the photocatalytic plasma, and which inhibit the production of PHPG for release into the environment as noted above.
In general, hydroxyl radicals are produced by the oxidation of water and require an oxidation potential of at least 2.85 eV to take place. The catalyst, therefore, must be activated by photons with at least this required energy. Photons with lower energy than 2.85 eV will not produce hydroxyl radicals, but photons with energy of at least 1.71 eV can photolyse hydrogen peroxide into hydroxyl radicals. Excess light with energy of 1.71 eV or above should be avoided due to the destruction of hydrogen peroxide.
Inside a plasma reactor, where it is possible for free electrons to recombine with hydroxyl radicals and form hydroxide ions, this is the thermodynamically favored reaction because it has the highest reduction potential, 2.02 eV. All reactions with lower reduction potentials, such as the combination of hydroxyl radicals to form hydrogen peroxide, 1.77 eV, are not favored. In rare instances where the formation of hydrogen peroxide occurs, a stoichiometric excess of two free electrons will be created. In this case the stoichiometric excess of free electrons makes it possible for lower potential reactions to take place, most notably the reduction of the hydrogen peroxide molecule into a hydroxyl radical and a hydroxide ion, 0.71 eV, then further down to water by separate reduction of the radical and of the ion.
In a plasma reactor, the abundance of free electrons ensures that the reduction of hydroxyl radicals dominates, and that any hydrogen peroxide that may theoretically be formed is immediately reduced back into water.
In contrast, in a PHPG reactor, production of hydrogen peroxide is favored because the reactor separates hydroxyl radicals from the free electrons, preventing the reduction of the hydroxyl radicals to water. This permits the next most favored reaction to take place, the combination of hydroxyl radicals to form hydrogen peroxide. The hydrogen peroxide can be reduced back down to water by decomposition (reaction of hydrogen peroxide molecules with each other), but this effect is minimized by ensuring that the hydrogen peroxide produced is dilute.
Also, since the PHPG reactor separates hydroxyl radicals from the free electron remaining on the catalyst, the free electrons are forced to reduce another species, in this case dioxygen. The reduction of dioxygen to the superoxide ion has a negative reduction potential, −0.13 eV, which indicates that it is non-spontaneous, but only slightly so. The non-spontaneity is overcome by the build-up of free electrons on the catalyst, creating an increasing thermodynamic reduction pressure. This non-spontaneous reaction is the first of four steps in the reduction of oxygen to hydrogen peroxide, the remaining three of which are all spontaneous. It is important to note that when all four of these steps are combined into a single reduction reaction, the overall potential is positive, or spontaneous. It is easy to overlook the fact that the non-spontaneous first step must take place in order for the three remaining spontaneous steps to follow.
The reduction of dioxygen to hydrogen peroxide, forced by the removal of hydroxyl radicals from the free electrons remaining on the catalyst, results in the desired production of yet more hydrogen peroxide, of course.
The reactions listed in Table 1 are the most pertinent. Other reactions, known in the art can be added and their relative contributions reactions on the catalyst surface determined by their relative potentials compared to the key reactions. Notably, as in the formation of ozone by plasma reactors, another high potential reaction is introduced that destroys hydrogen peroxide. To completely avoid ozone production, one need only avoid the use of light at wavelengths of 186 nm and below.
The wavelengths of light used to activate photocatalysts are also energetic enough to photolyse the peroxide bond in a hydrogen peroxide molecule and are also an inhibitor in the production of PHPG for release into the environment. Further, the practice of using wavelengths of light that produce ozone introduces yet another species into the photocatalytic plasma that destroys hydrogen peroxide.O3+H2O2⇄H2O to 2O2 
In practice, photocatalytic applications have focused on the production of a plasma, often containing ozone, used to oxidize organic contaminants and microbes. Such plasmas are primarily effective within the confines of the reactor itself, by nature have limited chemical stability beyond the confines of the reactor, and actively degrade the limited amounts of hydrogen peroxide gas that they may contain. Further, because the plasma is primarily effective within the reactor itself, many designs maximize residence time to facilitate more complete oxidation of organic contaminants and microbes as they pass through the reactor. Since hydrogen peroxide has such a high potential to be reduced, the maximized residence time results in minimized hydrogen peroxide output.
Also, most applications of photocatalysis produce environmentally objectionable chemical species. First among these is ozone itself, an intentional product of many systems. Further, since organic contaminants that pass through a reactor are seldom oxidized in one exposure, multiple air exchanges are necessary to achieve full oxidation to carbon dioxide and water. As incomplete oxidation occurs, a mixture of aldehydes, alcohols, carboxylic acids, ketones, and other partially oxidized organic species is produced by the reactor. Often, photocatalytic reactors can actually increase the overall concentration of organic contaminants in the air by fractioning large organic molecules into multiple small organic molecules such as formaldehyde.
Methods of vaporizing aqueous hydrogen peroxide solutions produce, at best, hydrated forms of hydrogen peroxide. Also, though photocatalytic systems are capable of producing hydrogen peroxide, they have multiple limitations that severely inhibit PHPG production for release into the environment. We have previously disclosed methods and devices for producing PHPG in U.S. application Ser. No. 12/187,755, published May 1, 2012, as U.S. Patent Publication No. 2009/0041617, and hereby incorporated by reference in its entirety.